Energy
Chap. 16
I. Definitions
I. Definitions
A. Energy
Energy is the ability to do work or produce
heat
I. Definitions
A. Energy
B. Heat
Heat is energy moving from one place to
another
Heat = Energy
For us, these terms will be used synonymously
I. Definitions
A. Energy
B. Heat
C. Temperature
A measure of the kinetic energy of the
particles in a substance
I. Definitions
A. Energy
B. Heat
C. Temperature
D. Endothermic
A description of a process that absorbs heat
I. Definitions
A. Energy
B. Heat
C. Temperature
D. Endothermic
E. Exothermic
A description of a process that gives off heat
II. Types of Energy
A. Kinetic
Energy of motion
II. Types of Energy
A. Kinetic
1. Mechanical
Moving objects
II. Types of Energy
A. Kinetic
1. Mechanical
2. Thermal
Heat energy (moving particles)
II. Types of Energy
A. Kinetic
B. Potential
Stored energy
II. Types of Energy
A. Kinetic
B. Potential
1. Gravitational
Energy that can be released as gravity acts
II. Types of Energy
A. Kinetic
B. Potential
1. Gravitational
2. Chemical
Energy stored in chemical bonds
II. Types of Energy
A. Kinetic
B. Potential
C. Radiant
Energy in the form of light
III. Measuring Heat (q)
III. Measuring Heat (q)
A. Units
III. Measuring Heat (q)
A. Units
1. joule
SI unit of energy (work). Work done by
applying one Newton force over one meter.
III. Measuring Heat (q)
A. Units
1. joule
2. calorie
Energy required to heat one gram of water by
1° C.
III. Measuring Heat (q)
A. Units
1. joule
2. calorie
3. Calorie
A nutritional calorie. 1 Cal = 1000 cal
III. Measuring Heat (q)
A. Units
1.
2.
3.
4.
joule
calorie
Calorie
kilocalorie
Equivalent to
calories
III. Measuring Heat (q)
A. Units
1.
2.
3.
4.
joule
calorie
Calorie
kilocalorie
Equivalent to 1000 calories
III. Measuring Heat (q)
A. Units
1.
2.
3.
4.
5.
joule
calorie
Calorie
kilocalorie
BTU
Energy required to heat 1 lb. water by 1º F.
Heat Unit Conversions
1 cal = 4.184 J
1000 cal = 1 Cal = 1 kcal
Self Check – Ex. 1
A reaction produces 3800 J
of heat. How many calories
is this?
Self Check – Ex. 2
A can of soda contains 150
Calories. How many joules
of energy is this?
III. Measuring Heat (q)
A. Units
B. Heat is related to temperature
III. Measuring Heat (q)
A. Units
B. Heat is related to temperature
Temp. = 35º C
Temp. = 65º C
Beaker #1
Which has more heat?
Beaker #2
III. Measuring Heat (q)
A. Units
B. Heat is related to temperature
C. Heat also depends on. . .
III. Measuring Heat (q)
A. Units
B. Heat is related to temperature
C. Heat also depends on. . .
1. Mass of material
Which beaker could melt more
ice (which has more heat)?
T2 = 85º C
T1 = 85º C
Beaker #1
Beaker #2
III. Measuring Heat (q)
A. Units
B. Heat is related to temperature
C. Heat also depends on. . .
1. Mass of material
2. Type of material
III. Measuring Heat (q)
A. Units
B. Heat is related to temperature
C. Heat also depends on. . .
D. Specific Heat
Amount of heat required to raise the
temperature of 1 gram of substance by 1º C
III. Measuring Heat (q)
A. Units
B. Heat is related to temperature
C. Heat also depends on. . .
D. Specific Heat
1. Some material takes a lot of
energy to raise its temperature
III. Measuring Heat (q)
A. Units
B. Heat is related to temperature
C. Heat also depends on. . .
D. Specific Heat
1. Some material takes a lot of
energy to raise its temperature
2. Some material takes less
III. Measuring Heat (q)
A. Units
B. Heat is related to temperature
C. Heat also depends on. . .
D. Specific Heat
1. Some material takes a lot of
energy to raise its temperature
2. Some material takes less
3. For water it’s 1 calorie/g ºC
Specific Heat Table
Substance
Water
Aluminum
Iron
Copper
Silver
Gold
Lead
Spec. Heat (c)
4.184 J/g ºC
0.89 J/g ºC
0.45 J/g ºC
0.387 J/g ºC
0.24 J/g ºC
0.129 J/g ºC
0.l28 J/g ºC
III. Measuring Heat (q)
A. Units
B. Heat is related to temperature
C. Heat also depends on. . .
D. Specific Heat
E. Calculation
q = m x c x ∆T
Self Check – Ex. 3
How much heat must be
applied to a 25 g chunk of
iron to raise its temperature
by 100ºC? (ciron = 0.45 J/g ºC)
Self Check – Ex. 4
How much heat must be
applied to a 25 g sample of
water to raise its temperature
by 100ºC? (ciron = 4.18 J/g ºC)
IV. Bond Energy
IV. Bond Energy
A. When bonds are broken
energy is
.
This is endothermic
IV. Bond Energy
A. When bonds are broken
energy is required. (positive)
This is endothermic
IV. Bond Energy
A. When bonds are broken
energy is required. (positive)
B. When bonds are formed
energy is
.
IV. Bond Energy
A. When bonds are broken
energy is required. (positive)
B. When bonds are formed
energy is released. (negative)
This is exothermic
IV. Bond Energy
A. When bonds are broken
energy is required. (positive)
B. When bonds are formed
energy is released. (negative)
C. The sum of the bond energies
gives an estimate of the
reaction energy
IV. Bond Energy
A. When bonds are broken
energy is required. (positive)
B. When bonds are formed
energy is released. (negative)
C. The sum of the bond energies
gives an estimate of the
reaction energy
1. Positive values = endothermic
IV. Bond Energy
A. When bonds are broken
energy is required. (positive)
B. When bonds are formed
energy is released. (negative)
C. The sum of the bond energies
gives an estimate of the
reaction energy
1. Positive values = endothermic
2. Negative values = exothermic
Self Check – Ex. 5
Draw Lewis structures for
each substance and calculate
the energy for the reaction
below.
2CO + O2
2CO2
V. Enthalpy Stoichiometry
V. Enthalpy Stoichiometry
A. Enthalpy represented by H
and enthalpy change by ∆H.
V. Enthalpy Stoichiometry
A. Enthalpy represented by H
and enthalpy change by ∆H.
B. Enthalpy change for a reaction
measured in kJ/mol.
V. Enthalpy Stoichiometry
A. Enthalpy represented by H
and enthalpy change by ∆H.
B. Enthalpy change for a reaction
measured in kJ/mol.
C. Exothermic reactions have
negative ∆H values. (+ ∆H for
endothermic)
V. Enthalpy Stoichiometry
A. Enthalpy represented by H
and enthalpy change by ∆H.
B. Enthalpy change for a reaction
measured in kJ/mol.
C. Exothermic reactions have
negative ∆H values. (+ ∆H for
endothermic)
D. Solving problems
Enthalpy Stoichiometry
Enthalpy Stoichiometry
1. Write a balanced equation
Enthalpy Stoichiometry
1. Write a balanced equation
2. Identify the units of the
unknown
Enthalpy Stoichiometry
1. Write a balanced equation
2. Identify the units of the
unknown
3. Write the ‘given’
Enthalpy Stoichiometry
1. Write a balanced equation
2. Identify the units of the
unknown
3. Write the ‘given’
4. Fill in conversion factors
V. Enthalpy Stoichiometry
A. Enthalpy represented by H
and enthalpy change by ∆H.
B. Enthalpy change for a reaction
measured in kJ/mol.
C. Exothermic reactions have
negative ∆H values. (+ ∆H for
endothermic)
D. Solving problems
1. Finding heat
Self Check – Ex. 6
How much heat is produced
when 75g of hydrogen is
burned in oxygen?
2H2 + O2
2H2O
∆H = - 484 kJ
V. Enthalpy Stoichiometry
A. Enthalpy represented by H
and enthalpy change by ∆H.
B. Enthalpy change for a reaction
measured in kJ/mol.
C. Exothermic reactions have
negative ∆H values. (+ ∆H for
endothermic)
D. Solving problems
1. Finding heat
2. Finding mass
Self Check – Ex. 7
What mass of hydrogen is
required to produce 8500 kJ of
energy?
2H2 + O2
2H2O
∆H = - 484 kJ
VI. Enthalpy Calorimetry
VI. Enthalpy Calorimetry
A. Calorimetry is the science of
heat measurement
VI. Enthalpy Calorimetry
A. Calorimetry is the science of
heat measurement
B. ΔH can be determined
experimentally if you measure:
1. Mass of water in calorimeter
VI. Enthalpy Calorimetry
A. Calorimetry is the science of
heat measurement
B. ΔH can be determined
experimentally if you measure:
1. Mass of water in calorimeter
2. Initial temperature
VI. Enthalpy Calorimetry
A. Calorimetry is the science of
heat measurement
B. ΔH can be determined
experimentally if you measure:
1. Mass of water in calorimeter
2. Initial temperature
3. Final temperature
VI. Enthalpy Calorimetry
A. Calorimetry is the science of
heat measurement
B. ΔH can be determined
experimentally if you measure:
1. Mass of water in calorimeter
2. Initial temperature
3. Final temperature
4. Mass of reactant used
Calculating ∆H for rxn
1. Determine the units for your
answer.
Calculating ∆H for rxn
1. Determine the units for your
answer.
2. Calculate the heat gained by the
water.
Q = mc∆T
cwater = 4.18 J/g ∙ ºC
Calculating ∆H for rxn
1. Determine the units for your
answer.
2. Calculate the heat gained by the
water.
3. Determine the heat lost by rxn.
Qgained = - Qlost
Calculating ∆H for rxn
1. Determine the units for your
answer.
2. Calculate the heat gained by the
water.
3. Determine the heat lost by rxn.
4. Convert grams reactant into
moles reactant.
Calculating ∆H for rxn
1. Determine the units for your
answer.
2. Calculate the heat gained by the
water.
3. Determine the heat lost by rxn.
4. Convert grams reactant into
moles reactant.
5. Divide heat lost (#3) by the
moles reactant (#4).
Self Check – Ex. 8
When a 10.0 g of NaOH is
added to 200 g of water the
temperature goes from 18.2ºC
to 31.6ºC. What is ∆H for the
reaction in kJ/mol?
VII. Hess’s Law
VII. Hess’s Law
A. When a series of equations are
added together, their enthalpy
changes are also added
Add these equations
N2 (g) + O2 (g)
2NO (g) + O2 (g)
2NO (g)
ΔH = +181 kJ
2NO2 (g) ΔH = -113 kJ
VII. Hess’s Law
A. When a series of equations are
added together, their enthalpy
changes are also added
B. Equations can be altered
VII. Hess’s Law
A. When a series of equations are
added together, their enthalpy
changes are also added
B. Equations can be altered
1. If an equation is reversed . . .
CO2 (g)
CO (g) + ½ O2 (g) ΔH = +283 kJ
VII. Hess’s Law
A. When a series of equations are
added together, their enthalpy
changes are also added
B. Equations can be altered
1. If an equation is reversed . . .
2. If the coefficients are multiplied
by a factor . . .
2C (g) + O2 (g)
CO (g)
ΔH = -111 kJ
* Using fractions is perfectly acceptable
X
½
Self Check – Ex. 9
ΔH = -572 kJ
2H2 (g) + O2 (g)
2H2O (l)
H2 (g) + O2 (g)
2H2O2 (l) ΔH = -188 kJ
2H2O2 (l)
2 H2O (l) + O2 (g) ΔH = ? kJ
Self Check – Ex. 10
2H2 (g) + O2 (g)
3O2 (g)
2H2O (g) ΔH = -483.6 kJ
2O3 (g)
3H2 (g) + O3 (g)
ΔH = +284 kJ
3H2O (g) ΔH = ? kJ
VII. Phase changes and
heat
The End